08-Sep-11

Chapter 2 Atoms, Molecules, and Ions

Dr. A. Al-Saadi

1

Preview Reviewing the fundamental chemical laws.  Understanding U d di the h modern d view i off atomic i structure, molecules, and ions.  Being familiar with main groups of elements in the periodic table.  Naming simple compounds, compounds ionic compounds, and getting the chemical formulas from the names. 

Dr. A. Al-Saadi

2

1

08-Sep-11

Chapter 2

Section 1

The Atomic Theory In 1808, John Dalton presented his theory. 1. An element is made up from extremely small particles called atoms. Atoms of a given element are identical, but are different for different elements. Dalton had no idea what an atom would look like! 2. Elements combine to form chemical compounds. 3. A chemical reaction eaction involves vo ves rearrangement ea a ge e t of o atoms; ato s; itt doesn’t create or destroy them. Atoms » Elements » Molecules (Compounds) Dr. A. Al-Saadi

3

Chapter 2

Section 1

The Atomic Theory

Combination of oxygen and carbon to form carbon dioxide

Dr. A. Al-Saadi

4

2

08-Sep-11

Chapter 2

Section 1

Law of Definite Proportion 

By Joseph Proust.

Different samples of a given compound always contain the same elements in the same mass ratio.

Dr. A. Al-Saadi

5

Chapter 2

Section 1

What Did Dalton Observe in CO Molecules? Mass of oxygen that combines with 1g of carbon

Ratio of mass of oxygen that combines with 1g of carbon

1 2 or 1 1

1.33g

2.66g

Dr. A. Al-Saadi

2

etc.

2 4 or etc. 1 1

6

3

08-Sep-11

Chapter 2

Section 1

Law of Multiple Proportions 

By John Dalton.

If two elements can combine to form more than one compound with each other, the masses of one element that combine with a fixed mass of the other element are in ratios of small whole numbers.

Dr. A. Al-Saadi

7

Chapter 2

Section 1

Law of Multiple Proportions 

Dr. A. Al-Saadi

For several compounds of nitrogen (N) and oxygen (O), the followingg results were observed:

Compound A

Mass of Nitrogen that combines with 1g of Oxygen 1.750 g

Compound B

0 8750 g 0.8750

Compound C

0.4375 g

8

4

08-Sep-11

Chapter 2

Section 1

Law of Multiple Proportions Mass of Nitrogen that combines with 1g of Oxygen

Mass of N in A

Compound A

1.750 g

Compound B

0.8750 g

Compound C

0.4375 g

=

Mass of N in C Mass of N in B Mass of N in C

4

A

B

C

1 =

2 1

The mass ratios shown can be readily described on basis of the ratios of number of atoms. Dr. A. Al-Saadi

9

Chapter 2

Section 1

Law of Conservation of Mass 

Matter can be neither created nor destroyed. Because matter is made up of atoms that are unchanged (masses and properties) in a chemical reaction, it follows that mass must be conserved as well.

Combination of oxygen and carbon to form carbon dioxide Dr. A. Al-Saadi

10

5

08-Sep-11

Chapter 2

Section 2

Early Imagination of the Atom 

What did Dalton think about the structure of an atom?     



Extremely small. Invisible. Has a mass. The smallest size ever of matter. No internal structure.

By mid 1800’s it became evident that atoms are divisible - there is an internal structure to the atom. (subatomic particles)

Dr. A. Al-Saadi

11

Chapter 2

Section 2

Cathode Ray Experiment

Dr. A. Al-Saadi

12

6

08-Sep-11

Chapter 2

Section 2

Cathode Ray Experiment 

The Deflection of the ray by a magnet indicates that the rayy is made upp of negatively g y charged g pparticles.

Thomson measured the charge-to-mass ratio as:

e  1.76  108 C/g m

Dr. A. Al-Saadi

13

Chapter 2

Section 2

Cathode Ray Experiment 

Dr. A. Al-Saadi

Cathode ray experiments revealed important conclusions about the structure off the h atom.  The ray is a stream of negatively charged particles (later on called electrons).  All atoms must contain electrons.  Since the atom is neutral overall, overall it must have a positively charged component.

Plum-pudding model suggested by Thomson

14

7

08-Sep-11

Chapter 2

Section 2

Mass of the Electron 

Mass of the electron:

Millikan’s Experiment (1917).

Oil d droplets l t

Dr. A. Al-Saadi

15

Chapter 2

Section 2

Radioactivity Types of spontaneous radioactive emission: 

 

α particles: have +ve charge and have mass that is 7300 time the mass of electron β particles: highspeed electrons. γ particles: highenergy light.

Dr. A. Al-Saadi

16

8

08-Sep-11

Chapter 2

Section 2

Rutherford Experiment

Dr. A. Al-Saadi

17

Chapter 2

Section 2

The Proton and the Nucleus 

Dr. A. Al-Saadi

Rutherford Experiment (1910)

18

9

08-Sep-11

Chapter 2

Section 2

The Nuclear Atom 

Rutherford’s Model vs. Thomson’s Model.

Thomson’s model (The plum-pudding model)

Rutherford’s model (The nuclear atom)

Dr. A. Al-Saadi

19

Chapter 2

Section 2

The Nuclear Atom Main components of atoms: Outside the nucleus:  Electrons: are responsible for the chemistry of the atom. Inside the nucleus:  Protons: are positively charged particles whose charge is equal i magnitude in it d to t that th t for f electrons. The simplest view of the atom Dr. A. Al-Saadi

20

10

08-Sep-11

Chapter 2

Section 2

The Nuclear Atom 







Nucleus is very tiny in terms of size. Each proton carries exactly the opposite charge of an electron. Almost all the atomic mass is concentrated in it (very dense)!! The mass of the proton is 1 67×10-24 g. 1.67×10 g If a nucleus were to have the size of a pea, it would weigh 250,000,000,000 kg!

10-8 cm = 100 picometer 1×10-12 1 picometer (pm) = ________m

Dr. A. Al-Saadi

21

Chapter 2

Section 2

The Nuclear Atom 



Dr. A. Al-Saadi

Rutherford’s model left one problem:  If H has a mass of 1, then He should have a mass of 2.  But its mass is 4! J. Chadwick (1932) discovered the neutrons; massive but uncharged particles.

22

11

08-Sep-11

Chapter 2

Section 3

Atomic Number and Mass Number

Mass number

23 11

Na

Atomic number Element symbol Element symbol (Na) = Sodium (Note that it is neutral) Mass number (A) = # of protons + # of neutrons Atomic number (Z) = # of protons For Na ion, the charge = # of protons – # of electrons Dr. A. Al-Saadi

23

Chapter 2

Section 3

Atomic Symbols in the Periodic Table Hydrogen (name) Sodium = [original name is Natriam]

Dr. A. Al-Saadi

24

12

08-Sep-11

Chapter 2

Section 3

Exercise

Atomic number = # of protons = 63 => Eu Atomic mass = # of protons + # neutrons = 63 + 88 Atomic charge = 63 – 60 = 3+ 3 The symbol is 151 63 Eu

Another exercise: For

53 26

Fe 2

# of protons = 26 # of neutrons = 53 – 26 Net charge = 2+ # of electrons = 26 – 2 = 24 Dr. A. Al-Saadi

25

Chapter 2

Section 3

Isotopes Isotopes show almost identical properties

Mass number

A Z

X

Atomic number

Ion (Cation)

Sodium-23 23 11 Dr. A. Al-Saadi

Na

Isotopes

Sodium-24

1

# of protons = 11

# of neutrons = 12

# of electrons = 10

26

13

08-Sep-11

Chapter 2

Section 3

Isotopes 



Isotopes of Hydrogen 

H d Hydrogen ((protium) i )



Deuterium



Tritium

The chemical properties of an element are determined by the electrons and protons, not the neutrons. Thus, isotopes are chemically alike.

Dr. A. Al-Saadi

27

Chapter 2

Section 3

How are Atomic Masses Measured?

The Mass Spectrometer Dr. A. Al-Saadi

28

14

08-Sep-11

Chapter 2

Section 3

Mass Spectrum of Xenon Three isotopes of neon exist in nature with different abundance.

Dr. A. Al-Saadi

29

Chapter 3

Section 3

Mass Spectrum of Carbon

1/10000 the size of large peak

Carbon exists in the form of three isotopes: 12C (98.93%) 13C (1.07%) 14C (< 0.001%).

Three isotopes of carbon are present in nature C-12 is the most abundant isotope of carbon. Dr. A. Al-Saadi

30

15

08-Sep-11

Chapter 2

Section 4

The Periodic Table 



Scientists noticed that chemical and physical properties p p of certain groups g p of elements are similar to one another. This led to the development of the periodic table.

Dr. A. Al-Saadi

31

Chapter 2

Section 4

The Modern Periodic Table

Ti

Dr. A. Al-Saadi

32

16

08-Sep-11

Chapter 2

Section 4

The Modern Periodic Table - Periods : horizontal rows - Families (Groups) : vertical columns Elements in the same family have similar chemical and physical properties

T i

- Arranged in order of increasing atomic number Dr. A. Al-Saadi

33

Chapter 2

Section 4

The Modern Periodic Table 





Metals: compose most of the periodic table. They have characteristic physical properties e.g.  High heat and electric conduction, malleability (hammered to sheet), ductility (pulled into wires).  Tend to lose electrons to form +ve ions. Fe2+ , Fe3+ , Na+ ,K+ , Ca2+ . Nonmetals: lack the physical properties of metals.  Tend to gain electrons to become –ve ions, like Cl-, F-, O2-, S2-.  Tend to bond with each other by forming (covalent bonds), such as Cl2, HCl,, N2O,, CO2 etc.  react with metals to form salt (ionic bonds), such as NaCl, CaF2, etc. Metalloids: have intermediate properties. 

Dr. A. Al-Saadi

Examples are B, Si, Ge, etc.

34

17

08-Sep-11

Chapter 2

Section 4

The Modern Periodic Table - The metallic properties increase as going from right to left across a period. - The groups in the periodic table are given special names. • Alkali metals.

T i

• Alkaline earth metals metals. • Chalcogens • Halogens. • Nobel gasses. • Transition metals.

Dr. A. Al-Saadi

35

Chapter 2

Section 4

The Modern Periodic Table -The The groups in the periodic table are given special names. • Alkali metals. • Alkaline earth metals. • Chalcogens • Halogens. • Nobel gasses. • Transition metals.

Dr. A. Al-Saadi

36

18

08-Sep-11

Chapter 2

Section 5

Atomic Masses 





Atomic mass is the mass of an atom expressed in atomic mass unit (amu). By definition used by modern systems, carbon-twelve 12C is assigned a mass of exactly 12 amu. O atomic One i mass unit i is i defined d fi d as tha mass exactly equal to 1/12 the mass of one carbon-12 atom.

Dr. A. Al-Saadi

37

Chapter 2

Section 5

Atomic Masses 



Carbon-12 (12 amu) provides the standard for measuring the atomic mass of the rest of elements. Example: Hydrogen atom 1H was found to be 8.3985% as massive as the C-12 atom. Can you find the atomic mass of a hydrogen atom in amu?

Mass 1H = 12 amu × 0.083985 = 1.0078 amu.

Dr. A. Al-Saadi

38

19

08-Sep-11

Chapter 2

Section 5

Average Atomic Masses Why the carbon in the periodic table has a mass p m of f 12.01 amu and not 12 amu??

Carbon exists naturally as a mixture of three isotopes, 12C, 13C and 14C and thus the atomic mass unit used for the carbon atom in i the h periodic i di table bl is i the h average value of the masses of those isotopes.

Dr. A. Al-Saadi

39

Chapter 2

Section 5

Average Atomic Masses The average atomic mass (or just the atomic mass) of the carbon atom = 98.89% of 12 amu (12C) + 1.11% of 13.0034 amu (13C) = (0.9889)(12 amu) + (0.0111)(13.0034 amu) = 12.011 amu 

Dr. A. Al-Saadi

That is applied for all the elements of the periodic table.

40

20

08-Sep-11

Chapter 2

Section 5

Average Atomic Masses 



Remember that, there is no a single carbon atom that has the mass of 12.010 amu. This is the average mass per carbon atom That is applied for all the elements of the periodic table. The mass of each element listed in the periodic table is an average value based on the isotopic composition of the naturally occurring element.

Dr. A. Al-Saadi

41

Chapter 3

Section 5

Exercise

Atomic mass = (mass of isotope 1 x fractional abundance of 1) + (mass of isotope 2 x fractional abundance of 2) + (mass of isotope 3 x fractional abundance of 3) + .....

Dr. A. Al-Saadi

42

21

08-Sep-11

Chapter 3

Section 5

Exercise

Atomic mass   mi xi i

187Re

is 62.60% with a mass of 186.956 amu.

Mass of Re = 186.207 amu = (186.956amu) (0.6260) + (?? amu) (0.3740) Answer is Mass of 185Re = 184.9533 amu = 185.0 amu Dr. A. Al-Saadi

43

Chapter 2

Section 6

Molecular Compounds & Ionic Compounds 



Dr. A. Al-Saadi

The force holds atoms together is called a chemical bond. bond Some types of chemical bonds are  Covalent bonds: Two atoms “usually nonmetals” can form a bond by sharing electrons to produce “molecular compounds”.  Ionic bonds: Two oppositely charged ions (a cation and an anion) “a metal and a nonmetal” can form a bond by attraction to produce “ionic compounds”. 44

22

08-Sep-11

Chapter 2

Section 6

Molecular Compounds 

Molecule : combination of at least two atoms in a specific arrangement held together by chemical bonds.    

May be an element or a compound. H2, hydrogen gas, gas is an element. element H2O, water, is a compound. They are also called “binary compounds”.

N2

CO

CH4

Dr. A. Al-Saadi

Chapter 2

45

Section 6

Molecular Compounds 

Diatomic Molecule: 

Homonuclear (2 of the same atoms) • Examples: H2, N2, O2, F2, Cl2, Br2, and I2



Polyatomic Molecule:    



Dr. A. Al-Saadi

Contain more than 2 atoms Most molecules May contain more than one element Examples: ozone, O3; white phosphorus, P4; water, H2O, and methane (CH4)

Heteronuclear (2 different atoms) • Examples: CO and HCl

46

23

08-Sep-11

Chapter 2

Section 6

Molecular Formulas 

Molecular formula: shows exact number of atoms of each element in a molecule.  



Subscripts indicate number of atoms of each element present in the formula. Example: H2O, NH3, C12H22O11 etc.

Structural St t l formula: f l shows h th the generall arrangement of atoms within the molecule.

Dr. A. Al-Saadi

47

Chapter 2

Section 6

Molecular Formulas Covalent-bonded Molecules

Dr. A. Al-Saadi

48

24

08-Sep-11

Chapter 2

Section 6

Molecular Formulas 

Allotrope: one of two or more distinct forms of an element.  

oxygen, O2 and ozone, O3 (allotropic forms of oxygen) diamond and graphite (allotropic forms of carbon)

Dr. A. Al-Saadi

49

Chapter 2

Section 6

Naming Molecular Compounds  

Only used for the n second element

Dr. A. Al-Saadi

It is also known as “Nomenclature” Nomenclature . Binary molecular (or covalent) compounds are composed of two nonmetals:  Name the first element.  Name the second element changing ending to “-ide”. ide .  If the two elements form more than one type of binary molecular compounds then use prefixes to indicate number of atoms of each element. 50

25

08-Sep-11

Chapter 2

Section 6

Naming Molecular Compounds

  

Only used for the n second element

   

HCl SiC NO N2O N2O5 SO2 PCl3

Hydrogen chloride Silicon carbide Nitrogen monoxide Dinitrogen monoxide Dinitrogen pentoxide Sulfur dioxide Phosphorus trichloride

Dr. A. Al-Saadi

51

Chapter 2

Section 6

Naming Molecular Compounds 

Name the following: 







Dr. A. Al-Saadi

Cl2O Dichlorine monoxide. CBr4 Carbon tetrabromide. ClO2 Chlorine dioxide. SO Sulfur monoxide.

52

26

08-Sep-11

Chapter 2

Section 6

Naming Molecular Compounds 

The names of molecular compounds containing hydrogen do not usually follow the systematic nomenclature guidelines.  B2H6 diborane  SiH4 silane  NH3 ammonia  PH3 phosphine h hi  H2O water  H2S hydrogen sulfide

Dr. A. Al-Saadi

53

Chapter 2

Section 6

Naming Binary Acids 



Dr. A. Al-Saadi

Acids when are dissolved in water, they produce H+ ions (protons) in the solutions solutions.  Examples are: HCl, HBr. Binary acids:  Many have 2 names • Pure substance: HCl, hydrogen chloride. • Aqueous solution: when dissolved in water it is called hydrochloric acid.

54

27

08-Sep-11

Chapter 2

Section 6

Naming Binary Acids 

In order to name binary acids:   



Remove the “–gen” –gen ending from hydrogen leaving “hydro–”. Change the “–ide” ending on the second element to “–ic”. Combine the two words and add the word “acid”.

Name the following: 



HBr Hydrogen bromide ; Hydrobromic acid H2S Hydrogen sulfide ; Hydrosulfuric acid

Dr. A. Al-Saadi

55

Chapter 2

Section 6

Naming Organic Compounds 

Organic Compounds: contain carbon and hydrogen y g (sometimes ( with oxygen, yg , nitrogen, g , sulfur and the halogens).  Hydrocarbons : contain only carbon and hydrogen.  Alkanes : simplest examples of hydrocarbons. Their names depend on the number of carbon atoms in the molecule. molecule



Inorganic Compounds: normally do not contain carbon.

Dr. A. Al-Saadi

56

28

08-Sep-11

Chapter 2

Section 6

Alkanes

Dr. A. Al-Saadi

57

Chapter 2

Section 6

Alkanes

Dr. A. Al-Saadi

58

29

08-Sep-11

Chapter 2

Section 6

Functional Groups 



Many derivatives of alkanes are derived by replacing a hydrogen with one or more functional groups. Functional group d determines i chemical h i l properties and is responsible for chemical reactions.

Dr. A. Al-Saadi

59

Chapter 2

Section 6

Empirical Formulas 

Empirical Formula: tells:  what elements are present in a molecule. molecule  In what whole-number ratio they are combined. Molecular(true)

H 2O 2 N 2H 4 H 2O Dr. A. Al-Saadi

Empirical(simplest)

HO O NH2 H 2O 60

30

08-Sep-11

Chapter 2

Section 6

Molecular and Empirical Formulas

Dr. A. Al-Saadi

61

Chapter 2

Section 7

Ionic Bonds and Ionic Compounds

11 protons 11 electrons

Na

17 protons 17 electrons

Cl Na+

-

Cl

e-

Na+ 11 protons 10 electrons

Dr. A. Al-Saadi

-

Cl

+ 17 protons 18 electrons

In form of crystals, called ionic solid or commonly known as salt 62

31

08-Sep-11

Chapter 2

Section 7

Ions and Ionic Compounds 

 



Ion: an atom or group of atoms that has a net ppositive or negative charge. Monatomic ion : one atom with a positive or negative charge. Cation : ion with a net positive charge due to the loss of one or more electrons. electrons Anion : ion with a net negative charge due the gain of one or more electrons.

Dr. A. Al-Saadi

Na+

-

Cl

63

Chapter 2

Section 7

Common Monoatomic Ions Cations: • Lithium ion (Li+) • Potassium ion ((K+) • Aluminum ion (Al3+)

Type I

Dr. A. Al-Saadi

• Iron (II) ion (Fe2+) • Iron (III) ion (Fe3+) ( 4+) • Lead ((IV)) (Pb • Lead (II) (Pb2+)

Type II

64

32

08-Sep-11

Chapter 2

Section 7

Common Monoatomic Ions Anions: • Fluoride ion (F–) • Oxide ion ((O2–) • Nitride ion (N3–)

Dr. A. Al-Saadi

65

Chapter 2

Type I

Section 7

Naming Binary Ionic Compounds +

-

It contains a +ve ion and a –ve ion.

1- Cations named first then anions. 22- Cation C ti element l t has h the th same name without change. 3-Use –ide root to the anion name. 4-Double check the ionic charges g to have the correct chemical formula. 5-You will need to practice this table.

Dr. A. Al-Saadi

6- You will need to be able to get names from formulas and vise versa.

66

33

08-Sep-11

Chapter 2

Type I

Section 7

Naming Binary Ionic Compounds Compound

Ions present

Name

NaCl

Na+ , Cl-

Sodium chloride

KI

K+ , I-

Potassium iodide

CaS

Ca2+ , S2-

Calcium sulfide

MgO

Mg2+ , O2-

Magnesium oxide

Al2O3

Al3+ , O2-

Aluminum oxide

Dr. A. Al-Saadi

67

Chapter 2

Section 7

Exercise      

Dr. A. Al-Saadi

Rb2O Rubidium oxide. CaS Calcium sulfide. AlI3 Aluminum iodide.

     

Strontium fluoride. SrF2 Aluminum selenide. Al2Se3 Magnesium phosphide. Mg3P2

68

34

08-Sep-11

Chapter 2

Type II

Section 7

Naming Binary Ionic Compounds 



Only for metals that can form more than one type yp off cations, the charge g must be specified using Roman numerals Examples:  CuCl Copper(I) chloride.  CuCl C Cl2 Copper(II) chloride.  CoCl3 Cobalt(III) chloride.

Dr. A. Al-Saadi

69

Chapter 2

Section 7

Polyatomic Ions 

Dr. A. Al-Saadi

Polyatomic ions : ions that are a combination of two or more atoms.

70

35

08-Sep-11

Chapter 2

Section 7

Polyatomic Ions

oxoanions

Dr. A. Al-Saadi

71

Chapter 2

Section 7

Exercise BaSO3 Barium sulfite  K2Cr2O7 Potassium dichromate  CuMnO4 C Copper(I) (I) permanganate t  NaNO2 Sodium nitrite 









Dr. A. Al-Saadi

Chromium(III) hydroxide Cr(OH)3  Magnesium cyanide Mg(CN)2  Lead(IV) carbonate Pb(CO3)2  Ammonium hypochlorite NH4ClO 









72

36

08-Sep-11

Chapter 2

Section 7

Naming Binary Compounds

Ionic compound Molecular compound

Dr. A. Al-Saadi

73

Chapter 2

Section 7

Naming Acids & Oxoacids 





Oxoacids : when ionized in water, they give H+ ions (protons)) and the corresponding (p p g polyatomic p y oxoanions in the solutions. Examples: HNO3, H2SO3, and HC2H3O2. When writing formulas, add the number of H+ ions necessary to balance the corresponding oxoanion’s negative charge. charge -

+

H

X

where X is an oxoanion +

H Dr. A. Al-Saadi

+

H

2-

X

74

37

08-Sep-11

Chapter 2

Section 7

Naming Acids & Oxoacids • HNO2. • HNO3. • H2SO3. • H3PO4.

• Hydrochloric y acid (HCl). ( ) • Hydrobromic acid (HBr).

• Nitric acid. • Phosphoric acid.

• Nitrous acid. • Sulfurous acid.

Dr. A. Al-Saadi

75

Chapter 2

Section 7

Naming Acids & Oxoacids

Dr. A. Al-Saadi

Acid

Anion

Name

HClO

Hypochlorite

Hypochlorous acid

HClO2

Chl it Chlorite

Chl Chlorous acid id

HClO3

Chlorate

Chloric acid

HClO4

Perchlorate

Perchloric acid

76

38

08-Sep-11

Chapter 2

Section 7

Hydrates 



Hydrates : compounds that have a specific number of water molecules within their solid structure  Hydrated compounds may be heated to remove the water forming an anhydrous compound Name the compound and add the word hydrate. Indicate the number of water molecules with a prefix on hydrate.  Example: CuSO4 · 5 H2O  Copper C (II) sulfate lf t pentahydrate t h d t

Dr. A. Al-Saadi

77

Chapter 2

Section 7

Exercise

Lead(II) acetate Copper(II) sulfate Calcium oxide Magnesium sulfate Magnesium hydroxide Calcium sulfate Dinitrogen monoxide

Dr. A. Al-Saadi

78

39